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This experiment aims to investigate the exothermic and endothermic reactions of various compounds using cold and hot packs within a calorimeter. Additionally, it seeks to analyze the relationship between the mass of salts and the heat absorbed or released during these reactions.
The invention of the first instant cold pack dates back to 1959 when Albert A. Robbins patented a device initially designed for temporary cooling of food and beverages. In 1971, pharmacist Jacob Spencer patented the first hot and cold pack, which was designed to comfortably conform to the body and could be reused (Rehak, 2014).
Hot and cold packs have since found diverse applications, with cold packs being effective in treating minor injuries by reducing pain and swelling, while hot packs are used to increase blood flow and facilitate the recovery of injured tissues.
Chemical reactions are classified as either exothermic or endothermic based on whether they release or absorb heat, respectively.
An exothermic reaction occurs when more energy is released during bond formation than is consumed during bond breaking.
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Conversely, an endothermic reaction consumes more energy during bond breaking than is released during bond formation. The heat of solution, expressed in kJ/mol, is associated with the enthalpy change when a solute dissolves in a solvent under constant pressure and infinite dilution (Mott).
Calorimetry plays a crucial role in temperature analysis during chemical reactions, as it allows for the measurement of heat transfer. A calorimeter is an isolated apparatus that prevents the exchange of energy with its surroundings. In experiments conducted under atmospheric pressure, any change in water temperature within the calorimeter is attributed to the enthalpy change resulting from the chemical reaction.
The fundamental equation for heat transfer is given by:
q = mcΔT
Where:
The heat capacity of a system determines the extent of temperature change for a given amount of energy transferred.
It is important to note that the heat of the chemical reaction, denoted as qrxn, is equal in magnitude but opposite in sign to the heat absorbed or released by the calorimeter (qcal). This relationship is expressed by the equation:
qrxn = -qcal
Here, ∆H represents the enthalpy change for the process, often associated with exothermic reactions such as combustion, and is typically measured in kJ/mol. To determine the ∆H for the reaction, it is necessary to multiply it by the number of moles involved to obtain the total heat released or absorbed.
The data collected during this experiment will enable us to determine the temperature change and whether the reaction is exothermic or endothermic. If the reaction results in heat release and a decrease in temperature, it is classified as exothermic. Conversely, if the reaction absorbs heat and leads to a temperature increase, it is categorized as endothermic. Furthermore, the data will aid in calculating the enthalpy change of the solution and exploring whether the mass of salts used correlates with the recorded temperature changes.
The experiment commenced by acquiring eight test tubes, each designated for a different salt sample. Five milliliters of diluted water were carefully added to each test tube. The test tubes were labeled with the names of their respective salts. Using a scoopula, approximately 2 cm of each salt was transferred into its corresponding test tube. The contents were thoroughly mixed with a stirring rod, and any temperature changes were observed by touching the test tubes to the palm of our hands. In cases where temperature changes were not easily detectable through touch, a temperature probe was employed. All temperature changes, categorized as either warm or cool, were meticulously recorded in our lab notebooks. This procedure was consistently replicated for all eight salts.
Following this initial assessment, we selected one salt from the hot pack and one from the cold pack for quantitative heat measurements. A modification was introduced at this stage, where instead of conducting a single trial, we conducted two trials for both the hot and cold packs, using two different masses (1g and 2g) for each pack.
We initiated the experiment with the hot pack, employing MgCl2 as the chosen salt. To carry out calorimetry, we utilized a styrofoam cup, measuring its initial mass. Subsequently, 25 mL of water were added to the cup, and the mass was measured again with the water. All measurements were diligently recorded in our lab notebook. The cup was secured in a ring stand to prevent it from falling while accommodating the temperature probe. The initial temperature of the water was recorded. Approximately 1g of the hot pack salt MgCl2 was precisely measured and added to the cup. The mixture was promptly swirled, and temperature changes were monitored using the temperature probe. Readings were noted when a constant temperature was achieved. The same procedure was repeated for the hot salt, this time using a different mass of approximately 2g.
Subsequently, the identical procedure was replicated for a salt from the cold pack, using both 1g and 2g masses.
| Salt | Chemical Formula | Temperature Change | Heat Released to Surroundings (Y/N) | Heat Absorbed from Surroundings (Y/N) | qsystem (< 0 or > 0) |
|---|---|---|---|---|---|
| Ammonium chloride | NH4Cl | Cooler | No | Yes | q < 0 |
| Magnesium sulphate | MgSO4 | Warmer | Yes | No | q > 0 |
| Calcium chloride | CaCl2 | Warmer | Yes | No | q > 0 |
| Magnesium chloride | MgCl2 | Warmer | Yes | No | q > 0 |
| Potassium nitrate | KNO3 | Cooler | No | Yes | q < 0 |
| Lithium chloride | LiCl | Warmer | Yes | No | q > 0 |
| Potassium chloride | KCl | Cooler | No | Yes | q < 0 |
| Sodium nitrate | NaNO3 | Cooler | No | Yes | q < 0 |
Table 12.2 displays the data obtained from the initial phase of the experiment, which aimed to categorize salts as either exothermic or endothermic reactions. A temperature change from normal to cooler indicates an endothermic reaction, while a change from normal to warmer signifies an exothermic reaction.
Based on the observations, the following conclusions can be drawn:
Moreover, the table also provides information on whether heat is released or absorbed from the system to its surroundings (qsystem). In cases where the temperature change is cooler, heat is absorbed (q < 0), signifying an endothermic reaction. Conversely, when the temperature change is warmer, heat is released (q > 0), indicating an exothermic reaction.
| Hot Pack Salt | Cold Pack Salt | ||||
|---|---|---|---|---|---|
| Trial 1 | Trial 2 | Trial 1 | Trial 2 | ||
| Parameters | Units | ||||
| Mass of empty cup | g | 2.290 | 2.293 | 2.283 | 2.288 |
| Mass of water | g | 26.287 | 26.045 | 28.838 | 24.534 |
| Mass of sample | g | 1.041 | 2.036 | 1.058 | 2.056 |
| Mass of solution (water+sample) | g | 27.034 | 28.08 | 29.90 | 26.59 |
| Initial temperature | °C | 20.4 | 21.2 | 21.3 | 22.0 |
| Final temperature | °C | 32.2 | 45.3 | 18.7 | 17.5 |
| ΔT (ΔT) | °C | 11.8 | 24.1 | -2.6 | -4.5 |
| Hsystem = qsystem = m (solution) * c * ΔT | Joules | 1348.8 | 2829.4 | -325.0 | -500.3 |
| Hsystem = qsystem = -qsurrounding | Joules | -1348.8 | -2829.4 | 325.0 | 500.3 |
| Molar mass of salt | g/mol | 48.33 | 48.33 | 101.11 | 101.11 |
| Moles of salt dissolved | mol | 0.0215 | 0.0421 | 0.0105 | 0.0204 |
| Hsystem = Hsystem / moles salt dissolved | Joules/mol | -62697.7 | -67258.9 | -30952.4 | -24524.5 |
| Uncertainty | 3.0 | 1.6 | 3.4 | 1.5 | |
| ΔT / gram of salt | °C/g | 11.3 | 11.8 | -2.5 | -2.2 |
| ΔT / $ | °C/$ | 393.3 | 803.3 | -28.9 | -50.0 |
Table 12.3 provides the data obtained during the hot and cold pack salt experiment, detailing two trials for each type of salt. The data is presented in a tabular format for clarity and comparison between trials and salt types.
The table includes the following parameters:
The data provided in this table will be utilized to calculate and analyze the heat transfer and enthalpy changes for the hot and cold pack salt reactions.
| Average T (°C) / Mass | Average T/$ | |
|---|---|---|
| The Cold Pack Salt | ||
| NH4Cl | -2.691 | -38.44 |
| MgSO4 | 4.846 | 121.15 |
| KNO3 | -2.323 | -25.81 |
| CaCl2 | ||
| KCl | -2.195 | -36.58 |
| MgCl2 | 11.586 | 386.2 |
| NaNO3 | -1.993 | -13.29 |
| LiCl | 5.11 | 56.78 |
| The Hot Pack Salt |
Table 12.4 presents the class average data for temperature per mass and the average cost per salt for both cold pack and hot pack salts. This data allows for a comparison between the salts in terms of their temperature per mass and cost-effectiveness.
The table includes the following information:
The data provides insights into the temperature changes per unit mass and cost for each salt, enabling a comparison of their effectiveness in hot and cold pack applications.
The objective of this experiment was to classify the reactions of various compounds as exothermic or endothermic and determine the influence of mass on the amount of heat absorbed or released. During the initial phase of the experiment, it was evident that the salts exhibited distinct exothermic and endothermic reactions upon dissolving them in water.
Salts that exhibited an exothermic reaction upon dissolution, characterized by the release of heat, included magnesium sulphate, calcium chloride, magnesium chloride, and lithium chloride. These salts are suitable for use in hot packs, as they generate heat when dissolved in water.
On the other hand, salts that demonstrated an endothermic reaction upon dissolution, resulting in a cooling sensation, included ammonium chloride, potassium nitrate, potassium chloride, and sodium nitrate. These salts are suitable for use in cold packs, as they absorb heat and produce a cooling effect when dissolved in water.
The experiment also revealed that the mass of the salts significantly affects the amount of heat released or gained. In the hot pack salt trials, when the mass of magnesium chloride increased from 1.041g (trial 1) to 2.036g (trial 2), the heat released in the second trial was greater, leading to a higher final temperature. Conversely, in the cold pack salt trials, when the mass of potassium nitrate increased from 1.058g (trial 1) to 2.056g (trial 2), the heat absorbed in the second trial was greater, resulting in a lower final temperature and a colder reaction.
Based on the results obtained in this experiment, magnesium chloride is recommended for use in hot packs due to its high average temperature change per gram of salt (T/mass), as indicated in Table 12.4. Magnesium chloride is highly soluble in water, and its dissolution results in an exothermic reaction that makes the solution hot.
For cold packs, ammonium chloride is the preferred choice, as it possesses the highest average temperature change per gram of salt (T/mass) among all the cold pack salts, as shown in Table 12.4. Ammonium chloride, an ionic compound, reacts with a non-ionic compound in water, creating a cold sensation. This reaction is endothermic, drawing energy from the surrounding environment and causing the solution to become cold.
It is important to acknowledge that errors may have occurred during measurements and salt mixing. Uncertainties could have arisen when measuring the masses of the salts. These factors may have contributed to variations in the experimental results.
In conclusion, this experiment successfully identified exothermic and endothermic reactions of salts and demonstrated the impact of salt mass on heat transfer. The results provide valuable insights for selecting salts for hot and cold packs, with magnesium chloride recommended for hot packs and ammonium chloride for cold packs.
Experiment on Exothermic and Endothermic Reactions with Calorimetry. (2024, Jan 14). Retrieved from https://studymoose.com/document/experiment-on-exothermic-and-endothermic-reactions-with-calorimetry
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